IIT JEE Chemistry - Study Guide - 8 - p-Block Elements

Sections in the chapter – Jauhar Text Book 12th Class



Study Plan

Group 13 elements

Day 1
8.1 Occurrence and their uses
8.2 General characteristics of Group 13 elements
8.3 Trends in chemical reactivity

Day 2
8.4 Aluminium: Extraction and properties
Practice problems: 8.1 to 8.8
Revision

Day 3

Group 14 Elements

8.5 Occurrence and uses
8.6 General characteristics of group 14 elements
8.7 Trends in chemical reactivity

Day 4
8.8 Forms of silica
8.9 Silicates
8.10 Silicones

Day 5

8.11 Tin and lead
P.P. 8.9 to 8.15
Revision

Day 6

Group 15 elements
8.12 Occurrence and uses
8.13 General characteristics of group 15 elements
8.14 Trends in chemical reactivity

Day 7

8.15 Production of phosphorus
8.16 Allotropic forms of phosphorus
8.17 Phosphine
8.18 Structure of some compounds of phosphorus

Day 8
P.P. 8.16 to 8.20

Group 16 elements

8.19 Occurrence and uses
8.20 General characteristics of group 16 elements
8.21 Trends in chemical reactivity

Day 9 and 10
8.22 Important compounds of group 16 elements
P.P 8.21 to 8.30

Day 11

8.23 Production of sulphur
8.24 allotropes of sulphur
8.25 Sulphuric acid

Group 17 elements

8.26 Occurrence and uses
8.27 General characteristics of group 17 elements

Day 12

8.28 Trends in chemical reactivity
8.29 Bleaching powder
8.30 Interhalogen compounds
P.P. 8.31 to 8.38

Day 13
Group 18 elements
8.31 Occurrence of noble gases
8.32 Isolation of noble gases and uses
8.32 General characteristics of group 18 elements
8.33 Compounds of noble gases

Day 14

pp. 8.39 to 8.41

Key factors to remember

Day 15
Revision Exercises: Very Short Answer questions 1 to 30

Day 16
Revision Exercises: Very Short Answer questions 31 to 50

Day 17
Short Answer Questions: 1 to 20

Day 18
Short Answer Questions: 21 to 40

Day 19
Short Answer Questions: 41 to 60

Day 20
Short Answer Questions: 61 to 85


Day 21
Competition File: Conceptual questions: 1 to 20

Day 22
Competition File: Conceptual questions: 21 to 47

Day 23
Competition File: Multiple choice questions: 1 to 15

Day 24
Competition File: Multiple choice questions: 16 to 30

Day 25
Competition File: Multiple choice questions: 31 to 43

Day 26
Fill in the blanks: 15

Day 27
True or False: 10

Day 28 to 30
Revision and test paper questions

IIT JEE Chemistry - Study Guide - 9. d and f -Block Elements

Sections in the chapter - Jauhar

Chemistry of d-Block elements

9.1 Definition and electronic configurations
9.2 General characteristics of transition elements
9.3 General trends in chemistry of transition elements
Practice problems 9.1 to 9.10
9.4 General trends in some important compounds
9.5 Occurrence and principles of extraction of iron
9.6 Occurrence and principles of extraction of copper
9.7 Occurrence and principles of extraction of silver
9.8 Occurrence and principles of extraction of zinc
9.9 Occurrence and principles of extraction of mercury
9.10 Steel and important alloys
P.P. 9.11 to 9.16
9.11 Some important compounds of transition metals
P.P. 9.17 to 9.22
9.12 Photography

Chemistry of f-Block elements

9.13 f-Block elements
9.14 General characteristics of Lanthanides and Actinides
9.15 Actinides
9.16 General characteristics of Actinides
9.17 Comparison of Actinide and Lanthanide series
P.P. 9.23 to 9.26

Key facts to remember
Conceptual Questions with Answers: 25
Revision Exercises
Very Short Answer questions: 35
Short Answer Questions: 55
Long Answer Questions: 24

Competition File
Typical questions 8
Objective Questions:
Muliple choice questions: 50
Fill in the blanks: 15
True or False: 10

Study Plan

10 days main study
10 days revision


Day 1
Chemistry of d-Block elements

9.1 Definition and electronic configurations
9.2 General characteristics of transition elements
9.3 General trends in chemistry of transition elements up to melting and boiling points

Day 2
9.3 General trends in chemistry of transition elements – remaining portion.
Practice problems 9.1 to 9.10

Day 3

9.4 General trends in some important compounds
9.5 Occurrence and principles of extraction of iron

Day 4
9.6 Occurrence and principles of extraction of copper
9.7 Occurrence and principles of extraction of silver
9.8 Occurrence and principles of extraction of zinc
9.9 Occurrence and principles of extraction of mercury
9.10 Steel and important alloys
P.P. 9.11 to 9.16

Day 5

9.11 Some important compounds of transition metals
P.P. 9.17 to 9.22
9.12 Photography

Day 6

Chemistry of f-Block elements

9.13 f-Block elements
9.14 General characteristics of Lanthanides and Actinides
9.15 Actinides
9.16 General characteristics of Actinides

Day 7

9.17 Comparison of Actinide and Lanthanide series
P.P. 9.23 to 9.26

Key facts to remember

Day 8

Conceptual Questions with Answers: 25

Day 9
Revision Exercises: Very Short Answer questions: 35

Day 10

Short Answer Questions: 1 to 30 55

Revision period

Day 11
Short Answer Questions: 31 to 55

Day 12
Competition File: Typical questions 8

Day 13

Competition File: Muliple choice questions: 1 to 25

Day 14
Competition File: Muliple choice questions: 26 to 50

Day 15

Fill in the blanks: 15

Day 16
True or False: 10

Day 17 to 20
Revision and test paper questions

Inorganic Chemistry - Study Guide - IIT JEE

Textbooks

NCERT Books for XI and XII

Jauhar for XI and XII

Arihant Prakashan

Inorganic Chemistry for Competitions by O P Tandon, G.R. Bathla & Sons, Meerut

Silicon

Molecular Formula Si

Silicon Si;

Atomic weight 28.0855
Atomic number 14;
valence 4
Melting point 1414 deg C;
Boiling point 3265 deg C;
sp. gr. 2.33 (25 deg C);

Berzelius, generally credited with the discovery, in 1824 succeeded in preparing amorphous silicon by removing the fluosilicates by repeated washings. Deville in 1854 first prepared crystalline silicon, the second allotropic form of the element.

Silicon is present in the sun and stars and is a principal component of a class of meteorites known as "aerolites". It is also a component of tektites, a natural glass of uncertain origin.

Natural silicon contains three isotopes.

Fourteen other radioactive isotopes are recognized.

Silicon makes up 25.7% of the earth's crust, by weight, and is the second most abundant element, being exceeded only by oxygen.

Silicon is not found free in nature, but occurs chiefly as the oxide and as silicates. Sand, quartz, rock crystal, amethyst, agate, flint, jasper, and opal are some of the forms in which the oxide appears. Granite, hornblende, asbestos, feldspar, claymica, etc. are but a few of the numerous silicate minerals.

Silicon is prepared commercially by heating silica and carbon in an electric furnace, using carbon electrodes. Several other methods can be used for preparing the element.

Amorphous silicon can be prepared as a brown powder, which can be easily melted or vaporized.

Crystalline silicon has a metallic luster and grayish color.

The Czochralski process is commonly used to produce single crystals of silicon used for solid-state or semiconductor devices. Hyperpure silicon can be prepared by the thermal decomposition of ultra-pure trichlorosilane in a hydrogen atmosphere, and by a vacuum float zone process.

This product can be doped with boron, gallium, phosphorus, or arsenic to produce silicon for use in transistors, solar cells, rectifiers, and other solid-state devices which are used extensively in the electronics and space-age industries.

Hydrogenated amorphous silicon has shown promise in producing economical cells for converting solar energy into electricity.

Silicon is a relatively inert element, but it is attacked by halogens and dilute alkali. Most acids except hydrofluoric, do not affect it.

Silicones are important products of silicon. They may be prepared by hydrolyzing a silicon organic chloride, such as dimethyl silicon chloride. Hydrolysis and condensation of various substituted chlorosilanes can be used to produce a very great number of polymeric products or silicones, ranging from liquids to hard, glasslike solids with many useful properties.

Elemental silicon transmits more than 95% bf all wavelengths of infrared, from 1.3 to 6.7 um.

Silicon is one of man's most useful elements. In the form of sand and clay it is used to make concrete and brick; it is a useful refractory material for high-tempemture work, and in the form of silicates it is used in making enamels, pottery, etc.

Silica, as sand, is a principal ingredient of glass, one of the most inexpensive of materials with excellent mechanical, optical, thermal, and electrical properties. Glass can be made in a very great variety of shapes, and is used as containers, window glass, insulators, and thousands of other uses.

Silicon tetrachloride can be used to iridize glass. Silicon is important in plant and animal life. Diatoms in both fresh and salt water extract silica from the water to build up their cell walls.

Silica is present in ashes of plants and in the human skeleton. Silicon is an important ingredient in steel;

silicon carbide is one of the most important abrasives and has been used in lasers to produce coherent light of 4560 A.

Miners, stonecutters, and other engaged in work where siliceous dust is breathed in large quantities often develop a serious lung disease known as silicosis.


http://www.speclab.com/elements/silicon.htm

IIT JEE Chemistry Ch.16A. Coordination Compounds

See for a set of questions on this topic the post

http://iit-jee-chemistry-ps.blogspot.com/2007/10/iit-jee-chemistry-questions.html
----------------------
Syllabus

Nomenclature of mononuclear coordination compounds, XII 10.1,2,3
Cis-trans and ionisation isomerisms, 10.4
Hybridization and geometries of mononuclear coordination compounds (linear, tetrahedral, square planar and octahedral).10.7

The topics are covered in detail Jauhar's XII book. The section numbers are given beside the topic.
-----------------------

Coordination compounds are a special class of compounds in which the centgral metal atom is surrounded by ions or molecules beyond their valency.

There are also referred to as coordination complexes or complexes.

Haemoglobin, Chlrophyll, and vitamin B-12 are coordinatio compounds of iron, magnesium and cobalt respectively.

The interesting thing of coordination compound is that these are formed from apparently saturated molecules capable of independent existence.

for example, when acqueous ammonia is addedt o green solution of nickel chloride, NiCl2, the colour changes to purple. The ni^2+ ions almost diappear from the solution. The solution on evaporation yields purple crystals corresponding to the formula [Ni(NH-3)-6]Cl-2. such a compound is called coordinatin compound. When this compound is now dissolved in water, there is hardly any evidence of Ni^2+ ions or NH-3 molecules. It ionizes to give a new species [Ni(NH-3)-6]^2+. the species in the square brackets does not ionise further. It remains as a single entity as an ion.

This is the unique feature of coordination compounds.

---------------
Nomenclature rules

From
http://www.iupac.org/publications/books/principles/principles_of_nomenclature.pdf

IUPAC booklet available for download at the above page id.

A coordination entity is composed of a central atom or atoms to which are attached
other atoms or groups of atoms, which are termed ligands. A central atom occupies a
central position within the coordination entity. The ligands attached to a central
atom define a coordination polyhedron. Each ligand is assumed to be at the vertex of
an appropriate polyhedron. The usual polyhedra are shown in Table 3.3 and they are
also listed in Table 4.4. Note that these are adequate to describe most simple
coordination compounds, but that real molecules do not always fall into these simple
categories. In the presentation of a coordination polyhedron graphically, the lines
defining the polyhedron edges are not indicative of bonds.

However, many ligands do not behave as donors of a single electron pair. Some
ligands donate two or more electron pairs to the same central atom from different
donor atoms. Such ligands are said to be chelating ligands, and they form chelate
rings, closed by the central atom. The phenomenon is termed chelation.

The number of electron pairs donated by a single ligand to a specific central atom
is termed the denticity. Ligands that donate one pair are monodentate, those that
donate two are didentate, those that donate three are tridentate, and so on.

Sometimes ligands with two or more potential donor sites bond to two (or more)
different central atoms rather than to one, forming a bridge between central atoms. It
may not be necessary for the ligand in such a system to be like ethane-l ,2-diamine,
with two distinct potential donor atoms. A donor atom with two or more pairs of
non-bonding electrons in its valence shell can also donate them to different centralatoms. Such ligands, of whatever type, are called bridging ligands. They bond to two
or more central atoms simultaneously. The number of central atoms in a single
coordination entity is denoted by the nuclearity: mononuclear, dinuclear, trinuclear,
etc. Atoms that can bridge include 5, 0 and Cl.

The original concepts of metal—ligand bonding were essentially related to the
dative covalent bond; the development of organometallic chemistry has revealed a
further way in which ligands can supply more than one electron pair to a central
atom. This is exemplified by the classical cases of bis(benzene)chromium and
bis(cyclopentadienyl)iron, trivial name ferrocene. These molecules are characterised
by the bonding of a formally unsaturated system (in the organic chemistry sense, but
expanded to include aromatic systems) to a central atom, usually a metal atom.

4.4.3 Mononuclear coordination compounds
4.4.3.1 Formulae. The central atom is listed first. The formally anionic ligands appear next,listed in alphabetical order of the first symbols of their individual formulae. The neutral ligands follow, also in alphabetical order. Polydentate ligands are included in alphabetical order, the formula to be presented as discussed in Chapter 3. The formula for the entire coordination entity, whether charged or not, is enclosed in square brackets. For coordination formulae, the nesting order of enclosing marks is as given on p. 13. The charge on an ion is indicated in the usual way by use of a right superscript. Oxidation states of particular atoms are indicated by an appropriate roman numeral as a right superscript to the symbol of the atom in question, and not in parentheses on the line. In the formula of a salt containing coordination entities, cation always precedes anion, no charges are indicated and there is no space between the formulae for cation and anion.

Examples
1. [Co(NH3)6]Cl
2. [PtC14]2

3. [CoC1(NH3)5]C1
6. [Cr"(NCS)4(NH3)2]
4. Na[PtBrC1(N02)(NH3)J
7. [Fe"(CO)4]2
5. [CaC12{OC(NH2)2}2]
The precise form of a formula should be dictated by the needs of the user.


The precise form of a formula should be dictated by the needs of the user. For
example, it is generally recommended that a ligand formula within a coordination
formula be written so that the donor atom comes first, e.g. [TiCl3(NCMe)3], but this
is not mandatory and should not affect the recommended order ofligand citation. It
may also be impossible to put all the donor atoms first, e.g. where two donors are
present in a chelate complex: [Co(NH2CH2CH2NH2)3]3. Whether the ethane-l,2-
diamine is displayed as shown, or simply aggregated as [Co(C2H8N2)3]3t is a matter
of choice. Certainly there is a conflict between this last form and the suggestion that
the donor atoms be written first. The aim should always be clarity, at the expense of
rigid adherence to recommendations.
It is often inconvenient to represent all the ligand formulae in detail. Abbreviations
are often used and are indeed encouraged, with certain provisos. These are: the
abbreviations should all be written in lower case (with minor exceptions, such as Me,
Et and Ph) and preferably not more than four letters; with certain exceptions of wide
currency, abbreviations should be defined in a text when they first appear; in a
formula, the abbreviation should be enclosed in parentheses, and its place in the
citation sequence should be determined by its formula, as discussed above; and
particular attention should be paid to the loss of hydrons from a ligand precursor.
This last proviso is exemplified as follows. Ethylenediaminetetraacetic acid
should be rendered H4edta. The ions derived from it, which are often ligands in
coordination entities, are then (H3edta), (H2edta)2, (Hedta)3 and (edta)4. This
avoids monstrosities such as edta-H2 and edtaH_2 which arise if the parent acid is
represented as edta. A list of recommended abbreviations is presented in Table 4.5.
4.4.3.2 Names. The addition of ligands to a central atom is paralleled in name construction.
The names of the ligands are added to that of the central atom. The ligands are listed
in alphabetical order regardless of ligand type. Numerical prefixes are ignored in this
ordering procedure, unless they are part of the ligand name. Charge number and
oxidation number are used as necessary in the usual way.
Of the two kinds of numerical prefix (see Table 4.2), the simple di-, tn-, tetra-, etc.
are generally recommended. The prefixes bis-, tris-, tetrakis-, etc. are to be used only
with more complex expressions and to avoid ambiguity. They normally require
parentheses around the name they qualify. The nesting order of enclosing marks is as
cited on p. 13. There is normally no elision in instances such as tetraammine and the
two adjacent letters 'a' are pronounced separately.

The names of ligands recommended for general purposes are given in Table 4.6.
The names for anionic ligands end in -o. If the anion name ends in -ite, -ate or -ide, the ligand name is changed to -ito, -ato or -ido. The halogenido names are, by
custom, abbreviated to halo. Note that hydrogen as a ligand is always regarded as
anionic, with the name hydride. The names of neutral and cationic ligands are never
modified. Water and ammonia molecules as ligands take the names aqua and
ammine, respectively. Parentheses are always placed around ligand names, which
themselves contain multiplicative prefixes, and are also used to ensure clarity, but
aqua, ammine, carbonyl (CO) and nitrosyl (NO) do not require them.
The names of all cationic and neutral entities end in the name of the element,
together with the charge (if appropriate) or the oxidation state (if desired). The
names of complex anions require modification, and this is achieved by adding the
termination -ate. All these recommendations are illustrated in the following examples.
----------------
Examples
1. Dichloro(diphenylphosphine)(thiourea)platinum(ii)
2. K4[Fe(CN)6]
3. [Co(NH3)6]C13
4. [CoCl(NH3)5}Cl2
5. [CoC1(N02)(NH3)4]Cl
6. [PtCl(NH2CH3)(NH3)2]Cl
7. [CuC12{OC(NH2)2}2]
8. K2[PdC14]
9. K[OsCl5N]
10. Na[PtBrC1(N02)(NH3)]
11. [Fe(CNCH3)6]Br2
12. [Ru(HSO3)2(NH3)4]
13. [Co(H20)2(NH3)4]C13
14. [PtC12(C5H5N)(NH3)]
15. Ba[BrF4]2
16. K[CrF4O]
17. [Ni(H20)2(NH3)4]S04
potassium hexacyanoferrate(ii)
potassium hexacyanoferrate(4—)
tetrapotassium hexacyanoferrate
hexaamminecobalt(iii) chloride
pentaamminechlorocobalt(2+) chloride
tetraamminechloronitrito-N-cobalt(iii) chloride
diamminechloro(methylamine)platinum(ii)
chloride
dichlorobis(urea)copper(ii)
potassium tetrachloropalladate(ii)
potassium pentachloronitridoosmate(2—)
sodium amminebromochloronitrito-
N-platinate( 1—)
hexakis(methyl isocyanide)iron(ii) bromide
tetraamminebis(hydrogensulfito)ruthenium(ii)
tetraamminediaquacobalt(iii) chloride
amminedichloro(pyridine)platinum(ii)
barium tetrafluorobromate(iii)
potassium tetrafiuorooxochromate(v)
tetraamminediaquanickel(ii) sulfate

***Table to be reformatted

-----------------------------

Designation of donor atom. In some cases, it may not be evident which atom in a
ligand is the donor. This is exemplified by the nitrito ligand in Examples 5 and 10,
p. 59. This can conceivably bind through an 0 or N atom. In simple cases, the donor
atom can be indicated by italicised element symbols placed after the specific ligand
name and separated from it by a hyphen, as demonstrated in those particular
examples. More complex examples will be dealt with below. With polydentate
ligands, this device may still be serviceable. Thus, dithiooxalate ion may be attached through S or 0, and formulations such as dithiooxalato-S,S' and dithiooxalato-0,0'should suffice. It could be necessary to use superscripts to the donor atom symbols if these need to be distinguished because there is more than one atom of the same kind to choose from.

Complicated examples are more easily dealt with using the kappa convention,
and this is particularly useful where a donor atom is part of a group that does not
carry a locant according to organic rules. The two oxygen atoms in a carboxylato
group demonstrate this. The designator i is a locant placed after that portion of the
ligand name that denotes the particular function in which the ligating atom is found.
The ligating atoms are represented by superscript numerals, letters or primes affixed
to the donor element symbols, which follow i without a space. A right superscript to
i denotes the number of identically bound ligating atoms.

Inclusion of structural information. The names described so far detail ligands and
central atoms, but give no information on stereochemistry. The coordination number
and shape of the coordination polyhedron may be denoted, if desired, by a
polyhedral symbol. These are listed in Table 4.4. Such a symbol is used as an affix in
parentheses, and immediately precedes the name, separated from it by a hyphen.
This device is not often used.
Geometrical descriptors, such as cis, trans, mer (from meridional) and fac (from
facial), have found wide usage in coordination nomenclature. The meaning is
unequivocal only in simple cases, particularly square planar for the first two and
octahedral for the others

More complex devices have been developed that are capable of dealing with all
cases. The reader is referred to the Nomenclature of Inorganic Chemistry, Chapter
10.
------------------------
Ionisation Isomers

Ionisation isomers:
Molecular structural formula is same. But different isomers give different ions in solution.

one isomer [PtBr(NH3)3]NO2 -> gives NO2- anions in solution
another isomer [Pt(NH3)3(NO2)]Br -> gives Br- anions in solution

Notice that both anions are necessary to balance the charge of the complex, and that they differ in that one ion is directly attached to the central metal but the other is not.

Geometric Isomers or Cis-Trans isomers

Geometric isomers are two or more coordination compounds which contain the same number and types of atoms, and bonds (i.e., the connectivity between atoms is the same), but which have different spatial arrangements of the atoms.

Not all coordination compounds have geometric isomers.

For example, in the square planar molecule, Pt(NH3)2Cl2, the two ammonia ligands (or the two chloride ligands) can be adjacent to one another or opposite one another.

Note that these two structures contain the same number and kinds of atoms and bonds but are non-superimposable. The isomer in which like ligands are adjacent to one another is called the cis isomer. The isomer in which like ligands are opposite one another is called the trans isomer.

For the common structures which contain two or more different ligands, geometric isomers are possible only with square planar and octahedral structures (i.e., geometric isomers cannot exist for linear and tetrahedral structures).

cis-[Co(NH3)4Cl2]+
Note that the two chloride ligands are adjacent to one another in this octahedral complex ion. In aqueous solution, this complex ion has a violet color.

trans-[Co(NH3)4Cl2]+
Note that the two chloride ligands are opposite one another in this complex ion. In aqueous solution, this complex ion has a green color.




------------------
Material about all isomers - In syllabus only ionic and cis-trans are specially mentioned

http://www.chem.purdue.edu/gchelp/cchem/whatis2.html

Coordination Isomers
Coordination isomers are two or more coordination compounds in which the composition within the coordination sphere (i.e., the metal atom plus the ligands that are bonded to it) is different (i.e., the connectivity between atoms is different).

Not all coordination compounds have coordination isomers.

Coordination isomers have different physical and chemical properties.

Example

[Cr(NH3)5(OSO3)]Br
Note that the sulfate group is bonded to the Cr atom (via an O atom) and is within the coordination sphere. Note also the octahedral structure. The bromide counterion is needed to maintain charge neutrality with the complex ion (i.e., [Cr(NH3)5(OSO3)]+) and is not shown in the structure.

[Cr(NH3)5Br]SO4
Note that the bromine atom is bonded to the Cr atom and is within the coordination sphere. Note also the octahedral structure. The sulfate counterion is not shown in the structure.


Linkage Isomers
Linkage isomers are two or more coordination compounds in which the donor atom of at least one of the ligands is different (i.e., the connectivity between atoms is different).

This type of isomerism can only exist when the compound contains a ligand that can bond to the metal atom in two (or more) different ways. Some ligands that can form linkage isomers are shown below.

Not all coordination compounds have linkage isomers.

Linkage isomers have different physical and chemical properties.

[Co(NH3)4(NO2)Cl]+
Note that the N atom of the nitrite group is bonded to the Co atom. The nitrite group is written as "NO2" in the molecular formula (rather than "ONO") with the N atom nearest to the Co symbol to indicate that the N atom (rather than an O atom) is the donor atom. Note also the octahedral structure.

[Co(NH3)4(ONO)Cl]+
Note that one of the O atoms of the nitrite group is bonded to the Co atom. The nitrite group is written as "ONO" in the molecular formula (rather than "NO2") with the O atom nearest to the Co symbol to indicate that the O atom is the donor atom. Note also the octahedral structure.

Geometric Isomers
Geometric isomers are two or more coordination compounds which contain the same number and types of atoms, and bonds (i.e., the connectivity between atoms is the same), but which have different spatial arrangements of the atoms.

Not all coordination compounds have geometric isomers.

For example, in the square planar molecule, Pt(NH3)2Cl2, the two ammonia ligands (or the two chloride ligands) can be adjacent to one another or opposite one another.

Note that these two structures contain the same number and kinds of atoms and bonds but are non-superimposable. The isomer in which like ligands are adjacent to one another is called the cis isomer. The isomer in which like ligands are opposite one another is called the trans isomer.

For the common structures which contain two or more different ligands, geometric isomers are possible only with square planar and octahedral structures (i.e., geometric isomers cannot exist for linear and tetrahedral structures).

cis-[Co(NH3)4Cl2]+
Note that the two chloride ligands are adjacent to one another in this octahedral complex ion. In aqueous solution, this complex ion has a violet color.

trans-[Co(NH3)4Cl2]+
Note that the two chloride ligands are opposite one another in this complex ion. In aqueous solution, this complex ion has a green color.

Optical Isomers
Optical isomers are two compounds which contain the same number and kinds of atoms, and bonds (i.e., the connectivity between atoms is the same), and different spatial arrangements of the atoms, but which have non-superimposable mirror images. Each non-superimposable mirror image structure is called an enantiomer. Molecules or ions that exist as optical isomers are called chiral.

Not all coordination compounds have optical isomers.

The Two Enantiomers of CHBrClF
Note that the molecule on the right is the reflection of the molecule on the left (through the mirror plane indicated by the black vertical line). These two structures are non-superimposable and are, therefore, different compounds.

Pure samples of enantiomers have identical physical properties (e.g., boiling point, density, freezing point). Chiral molecules and ions have different chemical properties only when they are in chiral environments.

Optical isomers get their name because the plane of plane-polarized light that is passed through a sample of a pure enantiomer is rotated. The plane is rotated in the opposite direction but with the same magnitude when plane-polarized light is passed through a pure sample containing the other enantiomer of a pair.



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web sites

http://www.chem.purdue.edu/gchelp/cchem/whatis2.html

Chapter 13A. Halogens

The Halogens are non-metals and form the 7th Group in the Periodic Table.

'Halogens' means 'salt formers' and the most common compound is sodium chloride which is found from natural evaporation as huge deposits of 'rock salt' or the even more abundant 'sea salt' in the seas and oceans.



Physical properties

• typical non-metals with relatively low melting points and boiling points.
• They are all poor conductors of heat and electricity - typical of non-metals.
• When solid they are brittle and crumbly e.g. iodine.
•
Colour
F Fluorine--- pale yellow gas

Cl Chlorine--- green gas

Br Bromine--- dark red liquid, brown vapour

I Iodine---- dark crumbly solid, purple vapour

At Astatine--- black solid, dark vapour

•
•


important trends down the Group with increasing atomic number
• The melting points and boiling increase steadily down the group (so the change in state at room temperature from gas ==> liquid ==> solid), this is because the weak electrical intermolecular attractive forces increase with increasing size of atom or molecule.
• They are all coloured non-metallic elements and the colour gets darker down the group.
• The size of the atom gets bigger as more inner electron shells are filled going down from one period to another.
Chemical Properties
• The atoms all have 7 outer electrons,
o they form singly charged negative ions e.g. chloride Cl- because they are one electron short of a noble gas electron structure. They gain one negative electron (reduction) to be stable and this gives a surplus electric charge of -1. These ions are called the halide ions, the bromide Br- and iodide I- ions.
o they form ionic compounds with metals e.g. sodium chloride Na+Cl-. (ionic bonding revision page)
o they form covalent compounds with non-metals and with themselves.
o The bonding in the molecule involves single covalent bonds e.g. hydrogen chloride HCl or H-Cl.
• The elements all exist as X2 or X-X, diatomic molecules where X represents the halogen atom.
• A more reactive halogen can displace a less reactive halogen from its salts .
• The reactivity decreases down the group .
• they are all TOXIC elements .
• Astatine is very radioactive, so difficult to study



Reaction with hydrogen H-2

• Halogens readily combine with hydrogen to form the hydrogen halides which are colourless gaseous covalent molecules. e.g. hydrogen + chlorine ==> hydrogen chloride
• H-2(g) + Cl-2(g) ==> 2HCl(g)
• The hydrogen halides dissolve in water to form very strong acids with solutions of pH1 e.g. hydrogen chloride forms hydrochloric acid in water HCl(aq) or H+Cl-(aq) because they are fully ionised in aqueous solution even though the original hydrogen halides were covalent. An acid is a substance that forms H+ ions in water.
• Bromine forms hydrogen bromide gas HBr(g), which dissolved in water forms hydrobromic acid HBr(aq). Iodine forms hydrogen iodide gas HI(g), which dissolved in water forms hydriodic acid HI(aq).

Reaction with Group 1 Alkali Metals Li, Na, K etc.

• Alkali metals burn very exothermically and vigorously when heated in chlorine to form colourless crystalline ionic salts e.g. NaCl or Na+Cl-.
• e.g. sodium + chlorine ==> sodium chloride
• 2Na(s) + Cl2(g) ==> 2NaCl(s)
• The sodium chloride is soluble in water to give a neutral solution pH 7, universal indicator is green. The salt is a typical ionic compound i.e. a brittle solid with a high melting point. Similarly potassium and bromine form potassium bromide KBr, or lithium and iodine form lithium iodide LiI.
•
Reaction with other metals
• If aluminium or iron is heated strongly in a stream of chlorine (or plunge the hot metal into a gas jar of chlorine carefully in a fume cupboard) the solid chloride is formed.
• aluminium + chlorine ==> aluminium chloride(white):
o 2Al(s) + 3Cl2(g) ==> 2AlCl3(s)
• iron + chlorine ==> iron(III) chloride(brown):
o 2Fe(s) + 3Cl2(g) ==> 2FeCl3(s)
• If the iron is repeated with bromine the reaction is less vigorous, with iodine there is little reaction, these also illustrate the halogen reactivity series.

JEE 2008 Syllabus Inorganic Chemistry

I am studying class XI and XII CBSE books of Jauhar along with TMH JEE Book of 2007. I am recording the sections and pages numbers where the specified topics are dealt with in these books for ready reference.

Such a reference seems to be essential for inorganic chemistry because the syllabus is not fitting into a simple chapter schemes of texts. For physical chemistry and organic chemistry each topic is distinct chapter in the books.



Inorganic Chemistry


Isolation/preparation and properties of the following non-metals:

Boron,XI Bk, Section 13.1, p.800-802
silicon, TMH JEE 2007 (TJ) page 13.4
nitrogen, XI Bk, Section 13.5,p.827-30
phosphorus, XII Bk, Sec 8.15
oxygen, XI, sec 13.7
sulphur and XII, 8.23,
halogens; XII, 8.26

Properties of allotropes of:
carbon (only diamond and graphite), XI, 812-814, XII, s 8.6.d.2, p 352
phosphorus and XII, 8.16, 369-70
sulphur.XII, 8.24, 389

Preparation and properties of the following compounds:

of sodium,

Oxides, XI, sec 12.6, 755-57, 763,764
peroxides, XI, 764
hydroxides, XI, 765
carbonates, XI, 767
bicarbonates,
chlorides and
sulphates


potassium,

Oxides,
peroxides,
hydroxides,
carbonates,
bicarbonates,
chlorides and
sulphates of


Compound of Sodium and Potassium are covered in unit 14 in TJ

magnesium

Oxides, XI, p774-76
peroxides,
hydroxides, XI 776
carbonates,
bicarbonates,
chlorides and XI 783
sulphates XI, 785


calcium;

Oxides, XI, 774, 786
peroxides,
hydroxides, XI, 787
carbonates,
bicarbonates,
chlorides and
sulphates XI, 779, 788


Compound of Magnesium and Calcium are covered in unit 14 of TJ

Boron:
diborane, XI, 13.2 806-10
boric acid and 804-5
borax; 802-4

Aluminium:
alumina,
aluminium chloride and
alums; - all in unit 14 TJ

Carbon:
oxides and XI , 13.4 818-21
oxyacid (carbonic acid);

Silicon:
silicones, XII, 8.10
silicates and 8.9
silicon carbide;

Nitrogen: XI 13.6
oxides, 833-36
oxyacids and 836-41
ammonia; 830-32

Phosphorus:
oxides, XII, p 367
oxyacids (phosphorus acid, phosphoric acid) and p 368, TJ p.15.17
phosphine; XII Sec 8.17, p370

Oxygen:
ozone and Xi sec 13.9, 848-853
hydrogen peroxide; XI sec 11.9, 729-738
XII, p 385

Sulphur:
hydrogen sulphide, XII p 364
oxides, XII sec 8.22.1 p 375
sulphurous acid, sec 8.22.2.1
sulphuric acid and 8.22.2.2, 8.25, p389
sodium thiosulphate; TJ p.15.21

Halogens:
hydrohalic acids, XII, 8.28.1 p 394
oxides and 394
oxyacids of chlorine, 95
bleaching powder; 8.29, 396

Xenon fluorides.XII, sec 8.33.A,

Transition elements (3d series):
Definition, XII, 9.1, 422
general characteristics, sec 9.2, 9.3
oxidation states and their stabilities, sec 9.3.6
colour (excluding the details of electronic transitions) and sec 9.3.7, Table 9.8 p430
calculation of spin-only magnetic moment; sec 9.3.9, p430

Preparation and properties of the following compounds:
Oxides and chlorides of tin and lead; XII, sec 8.11
Oxides, chlorides and sulphates of Fe2+, Cu2+ and Zn2+; XII sec 9.4, 9.11
Potassium permanganate, XII sec 9.11.6
potassium dichromate, 9.11.5
silver oxide,
silver nitrate, TJ p.14.9
silver thiosulphate.


Coordination compounds: XII, section 10
nomenclature of mononuclear coordination compounds, XII 10.1,2,3
cis-trans and ionisation isomerisms, 10.4
hybridization and geometries of mononuclear coordination compounds (linear, tetrahedral, square planar and octahedral).10.7


Ores and minerals: For this topic and topic of extraction metallurgy - refer Chapter X of Class XI book of Jauhar.

Commonly occurring ores and minerals of
iron,
copper,
tin,
lead,
magnesium,
aluminium,
zinc and
silver.

XI, sec 10.5


Extractive metallurgy: Chemical principles and reactions only (industrial details excluded); XI, 10.10
Carbon reduction method (iron and tin); 10.10.3.B.I.(i)
Self reduction method (copper and lead); 10.10.3.B.I.(vi)
Electrolytic reduction method (magnesium and aluminium); 10.10.3.B.III
Cyanide process (silver and gold). TJ p.17.6 (*gold is not there)



Principles of qualitative analysis: Ch. 18 of TMH JEE Book

Groups I to V (only Ag+, Hg2+, Cu2+, Pb2+, Bi3+, Fe3+, Cr3+, Al3+, Ca2+, Ba2+, Zn2+, Mn2+ and Mg2+);

Group I - TJp.18.3
Group II - 18.4
Group III- 18.5
Group IV - 18.5
Group V -- 18.6


Nitrate, 18.2
halides (excluding fluoride), 18.2
sulphate and 18.2
sulphide. 18.1


Except two or three topics all the topics are covered in the books. The specific items not covered can be accessed from one of the sites on the internet. I shall do it and post in the concerned chapter.

Study Guide Ch.13. NON-METALS

JEE Syllabus

Isolation/preparation and properties of the following non-metals:
Boron,
silicon,
nitrogen,
phosphorus,
oxygen,
sulphur and
halogens;

Properties of allotropes of
carbon (only diamond and graphite),
phosphorus and
sulphur.
---------------
MAIN TOPICS IN TMH BOOK

BORON
SILICON
NITROGEN
PHOSPHOROUS
OXYGEN
SULPHUR
HALOGENS

ALLOTROPES OF CARBON
------------
Boron(B)

Compounds relevant to study of Boron

Borax (Na two B four O seven) Calamnite (Ca two B six O eleven)
Boron Trioxide (B two O three) Boric Acid (H three B O three)
Boron Nitride (B N)

Sodium Carbonate (Na two C O three)



Boron (B) - Preparation and Properties

atomic number electronic configuration

Z = 5, 1s²2s²2px¹

main ores of Boron

Borax, Kernite, Colemanite and Orthoboric acid.

methods of obtaining Boron

1. By the reduction of boric oxide y an electropositive metal like magnesium.
2. By the reductin of volatile boron compounds by dihydrogen at high temperatures (1270K).
3. By the electrolytic reduction of fused borates or other boron compounds (e.g., KBF-4, potasium tetrafluoroborate) in molten KCL/KF at 1073K.
4. By the thermal decomposition of boron tri-iodided over red hot tungsten.
5. By thermal decomposition of boron hydrides and boron halides at about 1173K.

physical properties of Boron

1. Boron is an extremely hard solid next to diamond.
2. Its melting point is 2450K and boiling point is 3925K.
3. It is a poor conductor of heat electricity.
4. It has two isotopic forms - B-10 and B-11. Relative abundance 19% and 81% respectively.

chemical Properties of Boron

1. Combination with nonmetals: At room temperature, it reacts with flourine. Superficial reaction with oxygen. At higher termperatures, it reacts directly with all nonmetals except H, Ge, Te and nobles gases.
2. With water: It does not react with water even in the form of steam.
3. Acids: HCL does not react with boron.
When heated with concentrated sulphuric acid or nitric acid, boron is oxidized to boric acid.
Boron does not react with nonoxidizing acids.
4. Boron react with fused caustic alkalies like NaOH and KOH forming borates.
It dissolves in fused Na-2CO-3 and NaNO-3 mixture at 1173K.
5. Boron comines directly with almost all metals (except heavy metals) at higher temperatures.
Heavy Metals not combining with boron include Ag, Au, Cd, Hg, Ga, In, Tl, Pb, Sn, Bi etc.




Boron belongs to 13th group.
It has a very high melting point(2453 K).

It is extracted from minerals Borax(Na-2 B-4 O-7) or Calamnite (Ca-2 B-6 O-11)

Process of extracting Boron from Borax:
i) Borax is treated with concentrated hydrochloric acid. Boric acid is precipitated.
ii) boric acid is strongly heated. Boron trioxide is obtained (B-2 O-3).
iii) Boron trioxide is heated with Na, K or Mg pieces. Amorphous form of boron is obtained.

Alternatively mixture of boron tribromide vapours and hydrogen are passed over electrically heated filament of tungsten at 1200 degree centigrade. Crystalline form of Boron is obtained.

In another way, Boron trioxide can react with Aluminium to give Boron.



-------------------------
JEE question 2006

1. B(OH)-3 +NaOH two way reaction NaBO-2 +Na[B(OH)-4] + H-2O

How can this reaction is made to proceed in forward direction?
(A) addition of cis 1, 2 diol
(B) addition of borax
(C) addition of trans 1, 2 diol
(D) addition of Na-2HPO-4
Sol. (A)
-----------------



Silicon (Si)

Compounds relevant to study of silicon

Silica (Si O two)
Silicon Halide (Si X four) X is for halogen

silicon (Si) - Preparation and Properties

electronic configuration of Silicon

Atomic Number is 14. 1s²2s²2p^63s²3p²

methods of obtaining Silicon

1. Heating finely divided silica with magnesium powder.
2. Heating potassium silicoflouride with potassium metal.
3. Heating potassium silicoflouride wtih Al or Zn in an iron crucible.
4. Passing a current of SiCl-4 over molten aluminium.

physical Properties of Silicon

Silicon is available in two allotropic forms. the amorphous silicon and the crystalline or admantine silicon.

amorphous silicon is a dark brown powder which is insoluble in water.
Crystalline silicon forms pale yellow crystals.

chemical properties of silicon

1. Silicon burns in air or oxygen forming silicon dioxide.
2. With halogens, it forms halides, SiX-4.
3. With fused acqueous caustic alkalies, silicon forms alkali silicates with liberation of hydrogen.
4. silicon decomposes on red heating liberating hydrogen.
5. Metals like Magnesium and nonmetal carbon form silicides with silicon.
-------------------

Sulphur (S)

Compound relevant to study of Sulphur

Hydrogen Sulphide (H two S)
Sulphur Dioxide (S O two)
Sodium Thio-Sulphate (Na two S two O three)

Ions formed by Sulphur

Sulphide ion (S O three (two-))
thiosulphate ion (S two O three (two-))

Sulphur occurs in the native as well as combined form.Large quantites of sulphur are obtained from underground deposits in USA.

Partial combustion of Hydrogen sulphide produce sulphur.
Reaction of Hydrogen Sulphide with Sulphur dioxide also gives sulphur.
Sulphur occurs as S eight. It is a puckered ring with crown conformation.

several allotropic forms of sulphur are available.

Rhombic sulphur (or alpha sulphur) is the stablest form. It is obtained by evaporation of solution of sulphur in carbon disulphide.

At about 95 C - 96 C rhombic sulphur is changed into another allotropic form monoclinic sulphur (als known as prismatic or beta sulphur).
Monoclinic sulphure has needle shaped crystals.

Other allotropic forms are amorphous (or colloidal) and plastic sulphur (or Gamma sulphur).

Amorphous sulphur is obtained by (1) the action of dilute Hydrochloric acid on sodium thio sulphate solution, and
(2)by passing Hydrogen sulphide through dilute nitric acid.

Plastic sulphur is obtained by pouring boiling sulphur in cold water. This results in rapid cooling. Plastic sulphur consists of a completely random arrangement of chains of sulphur atoms. On standing it passes over to the crystalline rhombic sulphur.

Sulphur is also active element like Oxygen. It combines with a large number of metals and nonmetals.

Sulphur is oxidized by concentrated nitric acid and sulphuric acid.
It also reacts with hot concentrated solution of alkalies.

See for some more details and diagrams on sulphur

http://www.alevelchemistry.co.uk/Module_4/HTML%20Files/oxy_acids_of_non_metals/7.8.1_Sulphur_Chemistry_Notes.htm

Halogens

Flourine (F), Chlorine (Cl), Bromine (Br), Iodine (I).

Chlorine is yellow green gas
Bromine is reddish brown liquid
Iodine is steel grey solid

Allotrope

• Allotropes are elements that can exist in two or more different physical forms
• Diamond, graphite and buckminster fullerine are allotropes of carbon.
• The allotropes of carbon are all the element carbon. The type of carbon is determined from the bonding that occurs.

The carbon atoms in graphite are arranged in flat sheets that slide easily over each other, while the atoms in diamond are bonded in a complex, honeycombed structure that makes the solid much harder. The difference between the allotropes may be in the bonding present as in the diamond carbon example or it may be due to the crystal form produced as is common with most metals.
When compounds show a simlar property - different characteristic forms they are termed polymorphic
In 1830, Swedish chemist Jöns Berzelius (1779-1848) questioned whether a given element could exist in two or more forms with different chemical and physical properties. A decade later he suggested the name allotropy for this phenomenon. Berzelius knew of elements with this property--carbon, phosphorus, and sulfur in particular.

The properties of allotropes can be very different. Atmospheric oxygen (O2) and ozone (O3) are allotropes of oxygen. As a gas diatomic oxygen is essential for life whereas triatomic ozone is poisonous. Both forms occur naturally in the atmosphere, but diatomic oxygen is present throughout the atmosphere and ozone is normally present only in the upper atmosphere, except when it occurs as a pollutant at ground level or when it is manufactured by electrical discharge during storms. Diatomic oxygen is a more stable form than ozone.
With the exception of nitrogen, all group V elements show allotropy. For example, phosphorous occurs in three forms--white, red, and black. White phosphorous is poisonous and very reactive, red phosphorous is not poisonous and it is only moderately reactive, and black phosphorous is nearly inert. Each of these allotropes of phosphorous also has its own subset of forms.
Group VI elements also show allotropism. The classic example of this is rhombic and monoclinic sulfur. Both forms are ring structures containing eight sulfur molecules. Rhombic sulfur is the most stable form, but after heating and allowing the sulfur to melt and then cool down monoclinic sulfur is formed. Both are crystalline forms of sulfur, but their structures are slightly different. If monoclinic sulfur is left standing it will slowly revert back to the rhombic form over several days. These different forms of sulfur have slightly different characteristics, such as altered melting points. There are several other allotropes of sulfur that exist as well.
Sulfur also shows allotropes in the liquid form. These are generally due to the breaking down of the ring structure and having different ratios of broken and unbroken rings. This effect is also enhanced by unpaired electrons at the broken ends of the ring structure. These broken ends can bond with other broken chains giving long chain molecules with a very high viscosity and decreased mobility. This latter effect is due to increased tangling of the molecules. Sulfur vapor also shows allotropic behavior.


Allotropes of Carbon

Carbon has two allotropic forms - diamond and graphite. Other amorphous carbons are actually micro crystals of graphite.

Diamond, an allotrope of Carbon is face centered crystal. It is the hardest natural substance. It is a nonconductor of electricity. It has high refractive inded (2.45) and much of the light that falls on it is internally reflected.

Diamond burns in air at 900 C and in oxygen at 700 C and forms carbon dioxide.

Graphite has a layer like structure in the three dimensional space and this gives it the lubricating property.

The carbon content of various amorphous forms of graphite; Peat (60%), Lignite (70%), Bituminous coal (78%) Semibituminous coal (83%) and Anthracite Coal (90%). The residue that remains after destructive distilation of coal in the absence of air is coke.

Graphite
• Graphite is a slippery black powder.
• Graphite is the only nonmetallic substance that conducts electricity.
• Each C atom is bonded to 3 other C atoms, forming one double bond and 2 single bonds.
• Every C atom in the graphite structure is bonded in the same way, with each carbon atom having one double bond and 2 single bonds.
• Graphite consists of a two dimensional layer in which C atoms are arranged in a series of regular hexagons.
• There are weak attractions between layers and the layers can readily slide past one another.
• By experiment, it has been determined that the 3 bonds associated with the C atoms are identical. Thus, the fourth pair of electrons is really delocalized. This accounts for the electrical conductivity of graphite.

• Diamond is a colorless crystal.
• Diamond is the hardest naturally occurring substance.
• Each carbon atom is bonded to four other carbon atoms in a tetrahedral arrangement.
• The tetrahedral arrangement of carbon atoms gives diamond its characteristic
physical properties such as:

o diamond is very hard
o diamond has a high melting point
o diamond is not easily compressed
o when given a good cut, diamond reflects light off of its facets.
• Diamond is actually a giant interlocking group of tetrahedrally arranged atoms. The bond angle between any two bonds in the crystal is 109.5o.

Allotropes of Phosphorus

Phosphorus has three main allotropes: white, red and black.

White phosphorus is poisonous and can spontaneously ignite when it comes in contact with air. For this reason, white phosphorus must be stored under water and is usually used to produce phosphorus compounds.

Red phosphorus is formed by heating white phosphorus to 250°C (482°F) or by exposing white phosphorus to sunlight. Red phosphorus is not poisonous and is not as dangerous as white phosphorus, although frictional heating is enough to change it back to white phosphorus. Red phosphorus is used in safety matches, fireworks, smoke bombs and pesticides.

Black phosphorus is also formed by heating white phosphorus, but a mercury catalyst and a seed crystal of black phosphorus are required. Black phosphorus is the least reactive form of phosphorus and has no significant commercial uses.



Allotropes of sulphur

Sulfur exists as a number of different allotropes. Below 95.6°C, the stable crystal form is rhombic, while above this temperature the element changes to a triclinic form. Both these forms contain cyclic S8 molecules. At temperatures just above its melting point, sulfur is a yellow liquid also containing S8 molecules. At about 160°C, the sulfur atoms link together in chains and the liquid becomes dark brown and more viscuous. If the molten sulfur is quickly cooled, for example, by pouring it into cold water, the result is a reddish-brown solid called plastic sulfur. Above 200°C the viscosity decreases.

Sulfur vapor contains a mixture of S2, S4, S6, and S8 molecules. So-called "flowers of sulfur" is a yellow powder obtained by subliming the vapor.


Rhombic sulphur
Yellow transparent crystals
Melt pt. 113 oC
Obtained when sulphur crystallises
from solution.








Monoclinic sulphur
Amber coloured needles
Melt pt. 119 oC
Obtained when sulphur
solidifies above 95.6 oC.


At atmospheric pressure the rhombic form is stable below 95.6 oC and the monoclinic form above this temperature. Only at the transition temperature can the two allotropes exist in equilibrium with each other.


When sulphur is heated it melts and undergoes a series of changes as the temperature rises.



Ozone

Ozone (O three) is an allotropic form of Oxygen.
It is naturally formed above 20 Km from the earth from oxygen by absorbing sunlight. Ozone layer protects earth from concentration of ultraviolet rays. chlorofluorcarbons (C Cl two F two) used as refrigerant releases active chlorine after absorbing sunlight and this active chlorine decomposes ozone leading to destruction of ozone layer.







----------------------

JEE question

Statement - 1
Boron always forms covalent bond

Because
Statement - 2

The small size of B^3+favours formation of covalent bond.

(A) Statement – 1 is True, Statement – 2 is True; Statement – 2 is a correct explanation for statement – 1
(B) Statement – 1 is True, Statement – 2 is True; Statement – 2 is Not a correct explanation for Statement – 1.
(C) Statement – 1 is True, Statement – 2 is False
(D) Statement – 1 is False, Statement – 2 is True

answer A
This point is explicitly given in Jauhar's book
----------------------

JEE Question 2007 Paper I Linked Comprehension

The noble gases have closed-shell electronic configuration and are monoatomic gases under normal conditions. The low boiling points of the lighter noble gases are due to weak dispersion forces between the atoms and the absence of other interatomic interactions. The direct reaction of xenon with fluorine leads to a series of compounds with oxidation numbers +2, +4 and + 6. 4 XeF reacts violently with water to give XeO-3 . The compounds of xenon exhibit rich stereochemistry and their geometries can be deduced considering the total number of electron pairs in
the valence shell.


1. Argon is used in arc welding because of its

(A) low reactivity with metal
(B) ability to lower the melting point of metal
(C) flammability
(D) high calorific value

Solution: A

Argon provides inert atmosphere in welding due to low reactivity with metal.

2. The structure of XeO-3 is

(A) linear
(B) planar
(C) pyramidal
(D) T-shaped

Coorect choice: C

3. XeF-4 and XeF-6 are expected to be

(A) oxidizing
(B) reducing
(C) unreactive
(D) strong basic

Answer: A
----------------------------

Study Guide Ch.14. COMPOUNDS OF METALS

JEE Syllabus

Preparation and properties of the following compounds:
Oxides,
peroxides,
hydroxides,
carbonates,
bicarbonates,
chlorides and
sulphates of

sodium,
potassium,
magnesium and
calcium;

Aluminium: alumina, aluminium chloride and alums;
----------------

MAIN TOPICS IN TMH BOOK

SODIUM AND POTASSIUM
MAGNESIUM AND CALCIUM
ALUMINIUM
IRON
ZINC
---------------
Compounds of Sodium



Sodium oxide

Sodium oxide has formula Na2O.
It is also called sodium(I) oxide, disodium oxide, sodium monoxide, and soda.
It is used in ceramics as a glaze additive.
It is also a constituent of glass at around 15% sodium oxide


Sodium oxide has the formula weight of 61.979 u.

It is formed when sodium is burned with limited oxygen to the following equation:
4Na + O-2 → 2Na-2O

Sodium oxide is a basic compound, thus on reaction with water will create sodium hydroxide (NaOH).

Na-2O + H-2O → 2NaOH

sodium peroxide
A nearly white compound (Na 2 O 2 ), having vigorous oxidizing properties, and used in bleaching mechanical paper pulps and as a final stage in the bleaching of chemical paper pulps in some multi-stage bleaching sequences.

It is dangerous in use because, when in contact with organic matter, it reacts so vigorously with atmospheric moisture that sufficient heat can be generate to cause organic matter to burn or even explode.
sodium hydroxide
sodium hydroxide chemical compound, NaOH, is a white crystalline substance that readily absorbs carbon dioxide and moisture from the air.
It is very soluble in water, alcohol, and glycerin.

It is a caustic and a strong base
It is commonly known as caustic soda, lye, or sodium hydrate.
The principal method for its manufacture is electrolytic dissociation of sodium chloride; chlorine gas is a coproduct.
Small amounts of sodium hydroxide are produced by the soda-lime process in which a concentrated solution of sodium carbonate (soda) is reacted with calcium hydroxide (slaked lime); calcium carbonate precipitates, leaving a sodium hydroxide solution.
The major use of sodium hydroxide is as a chemical and in the manufacture of other chemicals; because it is inexpensive, it is widely used wherever a strong base is needed.
It is also used in producing rayon and other textiles, in making paper, in etching aluminum, in making soaps and detergents, and in a wide variety of other uses.

Sodium carbonate (Na2CO3)

Sodium carbonate exists as anhydrous (Na2CO3) and also as hydrated salt. The decahydrated salt (Na2CO3.10H2O) is known as washing soda while the anhydrous salt is called soda ash.

Occurrence

Large deposits of this salt occur in Owens lake in California and Lake Magadi in British East Africa. It occurs native as Na2CO3.NaHCO3.H2O in Egypt.

During hot weather, soda is also collected from a large number of alkaline lakes.

Manufacture of Sodium Carbonate

Ammonia-soda process (or Solvay process)

This process is the most popularly used method. As Ernest Solvay, the Belgian chemical engineer, devised it in 1864 it is known as Solvay process.

Raw materials

The raw materials for this process are common salt, ammonia and limestone (for supplying CO2 and quicklime).

Principle

When carbon dioxide is passed into a concentrated solution of brine saturated with ammonia, ammonium bicarbonate is produced,



The ammonium bicarbonate then reacts with common salt forming sodium bicarbonate,


Sodium bicarbonate being slightly soluble (in presence of sodium ions) gets precipitated. The precipitated sodium bicarbonate is removed by filtration and changed into sodium carbonate by heating.


The mother liquor remaining after the precipitation of sodium bicarbonate contains ammonium chloride. This is used to regenerate ammonia (one of the raw materials) by steam heating with milk of lime.


Lime is obtained by heating limestone.


Ammonia and carbon dioxide liberated are utilized in making the whole process cyclic and continuous. The only by-product in the process is calcium chloride.

Sodium Bicarbonate NaHCO-3
Sodium Bicarbonate, commonly called baking soda, is a white odourless, crystalline solid, completely soluble in water but slightly soluble in ethanol. It is the mildest of all sodium alkalis.

It is prepared from purified sodium carbonate or sodium hydroxide solution with passing carbon dioxide which is bubbled into the solution of pure carbonate, and the bicarbonate precipitates out to be dried as the bicarbonate is less soluble than the carbonate.

Sodium bicarbonate is also made as an intermediate product in the Solvay process (described above)which is to make sodium carbonate from calcium carbonate by treating sodium chloride with ammonia and carbon dioxide.

The major use of sodium bicarbonate is in baking powders.
Sodium Bicarbonate plays an important role in the products of many diverse industries with functions of releasing CO2 when heated above about 50 C or when reacted with a weak acid makes sodium bicarbonate a key ingredient in food leavening as well as in the manufacture of effervescent salts and beverages.
It can react as an acid or a base in water treatment.
In health and beauty applications, mild abrasivity and ability to reduce odors chemically by neutralizing the acid by-products of bacteria are utilized.
It is also used in treating wool and silk, fire extinguishers, pharmacy, leather, oredressing, metallurgy, in cleaning preparations and industrial & chemical processe.

Uses

food & food processing, beverages , pharmaceuticals , animal foodstuffs , household cleaning products , rubber & plastics foam blowing , fire extinguishers & explosion suppression , effluent & water treatment, flue gas treatment , oil drilling , industrial & chemical processes

Sodium chloride (NaCl) or common salt is an ionic crystal consisting of equal numbers of sodium and chlorine atoms and is an essential component in the human diet, being found in blood sweat and tears.
Occurrence
Sodium chloride is abundant and can be found naturally occurring. It can be found in the mineral halite (pure rock salt) as well as in mixed evaporates in salt lakes.
Sea water also contains 2.7% by weight salt and constitutes 80% of the dissolved minerals in sea water.
Production
Sodium chloride is mined or obtained from brine, when water is added to salt deposits.
Alternatively, it is obtained from sea water. This is commonly known as sea salt and constitutes most table salt. It also contains some impurities.

Sodium chloride:
• Has a cubic crystalline structure
• Is clear when pure, although may also appear white, grey or brownish, depending upon purity
• Is soluble in water
• Is slightly soluble in other liquids
• Is odourless
• Has a characteristic taste
• Molten sodium chloride is an electrical conductor

Symbol NaCl

Atomic Weight 58.44
Eutectic Composition 23.31% NaCl
Melting Point 801°C
Boiling Point 1465°C
Density 2.17g/cm3
Refractive Index 1.5442
Mohs Hardness 2.5
Co-Efficient of Thermal Expansion @ 0°C 40x10-6
Solubility g/100g H2O at 0°C 35.7


Sodium chloride is used for:
• Windows for analytical instruments
• De-icing
• Food and cooking
• High power lasers
• To produce chlorine and sodium
• Historically it has been used as a form of currency


SODIUM SULPHATE

FORMULA Na2SO4
MOL WT. 142.04
SYNONYMS Disodium monosulfate;

PHYSICAL AND CHEMICAL PROPERTIES
PHYSICAL STATE Hygroscopic white powder, Odorless
MELTING POINT 880 - 888 C
BOILING POINT 1100 C (Decomposes)
SPECIFIC GRAVITY 2.66 - 2.75
SOLUBILITY IN WATER Soluble
pH Aqueous solution is neutral

Sodium sulfate is a white, orthorhombic crystalline solid at room temperatures ( a monoclinic structure at > 100 C, a hexagonal structure at > 250C).

It is reduced to sodium sulfide at high temperature.
But sodium sulfate is a stable compound which does not decompose and does not react with oxidising or reducing agents at normal temperatures.

It is neutral (pH of 7) in water.
Sodium sulfate is most soluble in water at 32.4 C (49.7g/100 g).

Commercial major source of sodium sulfate is salt cake (impure sodium sulfate), a by-product of hydrochloric acid production from sodium chloride by treatment with sulfuric acid.

Sodium sulfate is obtained also as a byproduct of rayon production and sodium dichromate production. The decahydrate is known as Glauber's salt.

About half of the world's production is from the natural mineral form of the decahydrate (mirabilite).
Anhydrous sodium sulfate is found in nature as the mineral thenardite (Na2SO4).
Other sodium sulfate minerals are metasideronatrite Na4Fe2(SO4)4(OH)213H2O, krohnkite Na2Cu(SO4)212H2O, and schairerite Na3(SO4)(F,Cl).

Sodium sulfate is consumed in four major categories; powder detergents as a processing aid and as a filler, wood pulp processing for making kraft paper, textile dyeing processes as a levelling agent to penetrate evenly, and molten glass process to remove small air bubbles.

Sodium sulfate is employed also as a raw material for the production of various chemicals.
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Compounds of Potassium

potassium (I) oxide
• Formula as commonly written: K2O
Synonyms
• potassium (I) oxide
• potassium oxide
• dipotassium oxide
Physical properties
• Colour: yellowish white to grey
• Appearance: crystalline solid
• Melting point: >763°C; 350°C (decomposes
• Density: 2350 kg m-3
Potassium Peroxide
Identifications
• Synonyms/Related:
o Dipotassium peroxide
o K2O2


POTASSIUM HYDROXIDE


FORMULA KOH
MOL WT. 56.1
SYNONYMS Potassium hydrate; Caustic potash; Lye;
Potassium Hydroxide, commonly called caustic potash with formula KOH, is a caustic compound of strong alkaline chemical dissolving readily in water, giving off much heat and forming a caustic solution.

It is a white deliquescent solid in the form of pellets obtained by concentration of purified electrolytic potassium hydroxide solution with very low chloride content.

PHYSICAL AND CHEMICAL PROPERTIES
PHYSICAL STATE odorless white, deliquescent solid
MELTING POINT 360 C
BOLING POINT 1320 C
SPECIFIC GRAVITY 2.044
SOLUBILITY IN WATER Soluble
pH 13.5 (0.1 molar solution)


It reacts violently with acid and is corrosive in moist air toward metals such as zinc, aluminium, tin and lead forming a combustible, explosive gas.

It absorbs rapidly carbon dioxide and water from air.

Contact with moisture or water will generate heat.

Sodium hydroxide (Caustic soda) and potassium hydroxide (Caustic potash) are the two most important caustics. They are closely resembles in chemical properties and applications, e.g., in manufacturing liquid soap, in bleaching, and in manufacturing chemicals.

Potassium hydroxide is the largest-volume potassium chemical for non-fertilizer use.

Potassium Hydroxide is used in chemical manufacturing including potassium carbonate and other potassium chemicals, fertilizers, phosphates, agrochemicals, alkaline batteries and dyes.


APPLICATIONS
Potassium Hydroxide is used in chemical manufacturing including potassium carbonate and other potassium chemicals, fertilizers, phosphates, agrochemicals, alkaline batteries and dyes. It is also widely used in soap and bleaching industry.

Potassium carbonate

Potassium carbonate is a white salt, soluble in water (insoluble in alcohol), which forms a strongly alkaline solution. It can be made as the product of potassium hydroxide's absorbent reaction with carbon dioxide. It is deliquescent, often appearing a damp or wet solid. Potassium carbonate is used in the production of soap and glass.


Production

Potassium carbonate is prepared commercially by the electrolysis of potassium chloride. The resulting potassium hydroxide is then carbonated using carbon dioxide to form potassium carbonate, which is often used to produce other potassium compounds.
2KOH + CO2 → K2CO3 + H2O

Uses
Potassium carbonate has been used for soap, glass, and china production.
In the laboratory, it may be used as a mild drying agent where other drying agents such as calcium chloride may be incompatible. However, it is not suitable for acidic compounds.

Mixed with water it causes an exothermic reaction that results in a temperature change, producing heat.

In cuisine, it is used as an ingredient in the production of grass jelly, a food consumed in Chinese and Southeast Asian cuisines.

Potassium carbonate is being used as the electrolyte in many cold fusion experiments.

Potassium bicarbonate (also known as potassium hydrogen carbonate or potassium acid carbonate), is a colorless, odorless, slightly basic, salty substance. The compound is used as a source of carbon dioxide for leavening in baking, extinguishing fire in powder fire extinguishers, acting as a reagent, and a strong buffering agent in medications. The US Food and Drug Administration (FDA) recognizes potassium bicarbonate as "generally recognized as safe". It is used as a base in foods to regulate pH.
Potassium bicarbonate is soluble in water, and is often found added to bottled water to affect taste; however, it is not soluble in alcohol. Decomposition of the substance occurs between 100°C and 120°C into K2CO3 (potassium carbonate), H2O (water), and CO2 (carbon dioxide). In concentrations greater than 0.5%, KHCO3 can have phytotoxic effects on plants (potassium bicarbonate has widespread use in crops, especially for neutralizing acidic soil), although there is no evidence of human carcinogenicity, no adverse effects of overexposure, and no LD50.
Physically, potassium bicarbonate occurs as a crystal or a soft white granular powder. It has a CAS No [298-14-6]. It is manufactured by reacting potassium carbonate with carbon dioxide and water:
K2CO3 + CO2 + H2O → 2 KHCO3
Potassium bicarbonate is used as a fire suppression agent in some dry powder fire extinguishers,. It is about twice as effective in fire suppression as sodium bicarbonate.


potassium chloride (KCl)

Potassium chloride is also commonly known as "Muriate of Potash".
Potash varies in color from pink or red to white depending on the mining and recovery process used. White potash, sometimes referred to as soluble potash, is used primarily for making liquid starter fertilizers.
Manufacture/Extraction
Potassium chloride occurs naturally as sylvite, and in combination with sodium chloride as sylvinite.
It can be extracted from sylvinite.
It is also extracted from salt water .
It is a by-product of the making of nitric acid from potassium nitrate and hydrochloric acid.

Physical properties

In its pure state it is odorless.
It has a white or colorless vitreous crystal, with a crystal structure that cleaves easily in three directions.
Potassium chloride crystals are face-centered cubic.

Potassium chloride has a crystalline structure like many other salts. Structure: face-centered cubic. Lattice Constant: roughly 6.3 angstroms.
In chemistry and physics it is a very commonly used as a standard, for example as a calibration standard solution in measuring electrical conductivity of (ionic) solutions, since carefully prepared KCl solutions have well-reproducible and well-repeatable measurable properties.

Chemical properties

Potassium chloride can react as a source of chloride ion. As with any other soluble ionic chloride, it will precipitate insoluble chloride salts when added to a solution of an appropriate metal ion:
KCl(aq) + AgNO3(aq) → AgCl(s) + KNO3(aq)

Although potassium is more electropositive than sodium, KCl can be reduced to the metal by reaction with metallic sodium at 850 °C because the potassium is removed by distillation
(KCl(l) + Na(l) ⇌ NaCl(l) + K(g)
This method is the main method for producing metallic potassium.

As with other compounds containing potassium, KCl in powdered form gives a lilac flame test result.

KCl is used in medicine, scientific applications, food processing and in judicial execution through lethal injection.

Potassium sulfate (K2SO4)

Potassium sulfate (K2SO4) (also known as potash of sulfur) is a non-flammable white crystalline salt which is soluble in water. The chemical is commonly used in fertilizers, providing both potassium and sulfur.

Manufacture
Potassium sulfate can be synthesised by the decomposition of potassium chloride with sulfuric acid. Hhydrogen chloride evaporates and can be used to produce hydrochloric acid.

The Hargreaves method is basically the same process with different starting materials. Sulfur dioxide, oxygen and water (the starting materials for sulfuric acid) are reacted with potassium chloride. Hydrochloric acid evaporates off.

It is obtained as a by-product in many chemical reactions including the production of nitric acid. This can be done by mixing the following: 2 Parts Potassium Nitrate to 1 Part Sulfuric Acid (molar ratio).
2KNO3 + H2SO4 ---> 2HNO3 + K2SO4
To purify the crude product, it can be dissolved in hot water and then filtered and cooled, causing the bulk of the dissolved salt to crystallize with characteristic promptitude.


Properties
The anhydrous crystals form a double six-sided pyramid, but are in fact classified as rhombic.

They are transparent, very hard and have a bitter, salty taste.

The salt is soluble in water, but insoluble in solutions of potassium hydroxide (sp. gr. 1.35), or in absolute ethanol.

It melts at 1078 °C.


Uses
The principal use of potassium sulfate is as a fertilizer. The crude salt is also used occasionally in the manufacture of glass.









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Compounds of Magnesium

Magnesium oxide

Roasting either magnesium carbonate or magnesium hydroxide produces the oxygen compound magnesium oxide, commonly called magnesia, MgO, a white solid used in the manufacture of high-temperature refractory bricks, electrical and thermal insulators, cements, fertilizer, rubber, and plastics. It is used medically as a laxative.

Magnesium peroxide

Magnesium peroxide is a fine powder peroxide with a white to white-off color. It releases oxygen by breaking down at a controlled rate with a hydrous fluid.

In contact with water it decomposes by the reactions:

MgO2+ 2 H2O → Mg(OH)2 + H2O2
2 H2O2 → 2 H2O + O2

Applications
Magnesium peroxide being environmentally benign and its stable oxygen release are used widely in the cosmetic, agricultural, pharmaceutical, and environmental industries. It is used to reduce contaminant levels in groundwater. Magnesium peroxide is used in the bioremediation of contaminated soil and can improve the soil quality for plant growth and metabolism. It also used in the aquaculture industry for bioremediation.

Commercially, magnesium peroxide exists as a form of compound of magnesium peroxide and magnesium hydroxide

http://en.wikipedia.org/wiki/Magnesium_peroxide


Magnesium hydroxide

Magnesium hydroxide, Mg(OH)2, is a white powder produced in large quantities from seawater by the addition of milk of lime (calcium hydroxide). It is the primary raw material in the production of magnesium metal. In water it forms a suspension known as milk of magnesia, which has long been used as an antacid and a laxative.

Magnesium carbonate

Magnesium carbonate, MgCO3, occurs in nature as the mineral magnesite and is an important source of elemental magnesium. It can be produced artificially by the action of carbon dioxide on a variety of magnesium compounds. The odourless white powder has many industrial uses—e.g., as a heat insulator for boilers and pipes and as an additive in food, pharmaceuticals, cosmetics, rubbers, inks, and glass.


Magnesium chloride,

The action of hydrochloric acid on magnesium hydroxide produces magnesium chloride, MgCl2, a colourless, deliquescent (water-absorbing) substance employed in magnesium metal production, in the manufacture of a cement for heavy-duty flooring, and as an additive in textile manufacture.


Magnesium sulfate


Magnesium sulfate, MgSO-4, is a colourless, crystalline substance formed by the reaction of magnesium hydroxide with sulfur dioxide and air. A hydrate form of magnesium sulfate called kieserite, MgSO-4.H2O, occurs as a mineral deposit.

Synthetically prepared magnesium sulfate is sold as Epsom salt, MgSO-4.7H2O. In industry magnesium sulfate is used in the manufacture of cements and fertilizers and in tanning and dyeing; in medicine it serves as a purgative.


http://www.britannica.com/eb/article-4446/magnesium
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Compounds of Calcium

A number of calcium compounds such as calcium oxide, calcium hydroxide, calcium carbide, calcium chloride, calcium carbonate, calcium sulphate, etc. occur either in nature or are easily prepared. All these compounds find extensive use in our day-to-day living.

Calcium oxide (CaO)

Calcium oxide is commonly known as quicklime, and is a material of primary importance in the building industry.

Preparation

Quicklime (CaO) is prepared by heating very strongly, limestone (CaCO3) in a lime kiln.
Smaller pieces of limestone are introduced from the top and heating is done from the lower end.
Limestone decomposes at about 1000°C to give calcium oxide.


Properties

Calcium oxide (quicklime) is a white amorphous solid. On heating, quicklime (CaO) glows at high temperatures. This glow of white dazzling light is called lime light. Quicklime melts at 2870 K (2597°C).

On exposure to the atmosphere, it absorbs moisture and carbon dioxide to finally give calcium carbonate.


When water is poured over quicklime, a lot of heat is produced giving out steam with a hissing sound. This is called slaking of lime.


Uses

For the manufacture of bleaching powder, glass, calcium carbide, soda ash etc.

For white washing of buildings.

For tanning of leather.

As a fertilizer for acidic soil.

For softening hard water.

In building and construction industry as an important raw material.


Calcium peroxide (CaO2)

Calcium peroxide (CaO2) is a solid peroxide with a white or yellowish color.

When in contact with water it will immediately begin to decompose releasing oxygen.

It will dissolve in acid to form hydrogen peroxide.


Calcium hydroxide (Ca(OH)-2)

Calcium hydroxide in solid powdered form is called slaked lime. A suspension of slaked lime in water is called milk of lime.

A clear saturated solution of Ca(OH)2 is called lime water.

Preparation

From quicklime

Calcium hydroxide is obtained by treating quicklime with water


From calcium chloride

Calcium hydroxide is obtained as white precipitate when sodium hydroxide is added to a concentrated solution of calcium chloride.


Properties

Calcium hydroxide is a white amorphous solid (density: 2.08 g/mL).


Calcium hydroxide is sparingly soluble in water. The solubility decreases with temperature.

With carbon dioxide

When carbon dioxide is passed through lime water, it becomes milky due to formation of calcium carbonate.


The milkiness disappears on passing CO2 gas in excess because calcium carbonate changes to the soluble calcium bicarbonate.


When slaked lime is exposed to air, it absorbs carbon dioxide to form calcium carbonate.

This reaction forms the basis of the white washing of buildings.

With acids

Slaked lime (Ca(OH)2) dissolves in dilute hydrochloric acid.


However, it is not soluble in dilute sulphuric acid because the calcium sulphate formed is a water-insoluble salt.


With chlorine

Slaked lime reacts with chlorine to form bleaching powder.


Uses

As a building material.

For white washing buildings.

For the softening of hard water.
As lime water, it is used for the detection of carbon dioxide.

In tanning industry.





Calcium carbonate (CaCO3)

Occurrence

The most abundant mineral of calcium is calcium carbonate. It occurs in nature in different forms, such as limestone, marble, chalk etc.

Calcium carbonate occurs abundantly as dolomite, MgCO3.CaCO3, a mixture of calcium and magnesium carbonates. It is the chief constituent of shells of sea animals and also of bones along with tricalcium phosphate.

Preparation
Laboratory preparation

Calcium carbonate is prepared in the laboratory by passing carbon dioxide gas into lime water.


Calcium carbonate is also obtained by adding the solution of a soluble carbonate to soluble calcium salt e.g.,


The resulting precipitate is filtered, washed and dried. The product obtained is known as precipitated chalk.

Excess of carbon dioxide should be avoided since this leads to the formation of calcium hydrogen carbonate.




Properties

Calcium carbonate is a white fluffy powder. It is almost insoluble in water.

Action of heat

When heated to 1200 K, it decomposes to give quicklime and carbondioxide.
Action of heat

When heated to 1200 K, it decomposes to give quicklime and carbondioxide.


With acids

Calcium carbonate reacts with dilute acids to liberate carbon dioxide.


Uses

As a building material in the form of marble.


In the manufacture of quick lime.


As a raw material for the manufacture of sodium carbonate in Solvay process.


In the extraction of metals such as iron (as flux).


As a constituent of toothpaste.


Calcium sulphate or Gypsum (CaSO4)

Occurrence

In nature calcium sulphate occurs as

Anhydrite, CaSO4


Gypsum, CaSO4.2H2O


Large quantities of gypsum are available in India in the states of Punjab and Rajasthan.


Preparation


In laboratory, calcium sulphate is prepared by the action of sulphuric acid on calcium chloride, calcium oxide or calcium carbonate.


Calcium sulphate is obtained as dihydrate (CaSO4.2H2O) from the solution.

From gypsum

Calcium sulphate can be obtained by heating gypsum above 200°C.






Properties

Calcium sulphate exists as a white solid in two forms:

Dihydrate (CaSO4.2H2O)


Anhydrous (CaSO4)


It is sparingly soluble in water. The solubility increases upto 40°C, beyond which the solubility decreases.

Action of heat

During initial heating of calcium sulphate dihydrate, the crystal structure

changes. On further heating at 390 K, it loses water and forms CaSO4.5H2O or (CaSO4)2.H2O. This is known at Plaster of Paris. On heating above 437 K, it becomes anhydrous CaSO4 and does not set when mixed with water. It is known as Dead Burnt Plaster.

Above 1475 K (< 1200°C) it decomposes to give CaO.




With calcium sulphide

When heated with calcium sulphide at 1475 K (1200°C), it gives calcium oxide, and sulphur dioxide.




Uses

For making plaster of paris and ammonium sulphate.


Gypsum (CaSO)4.2H2O) is used as a fertilizer as well as in the manufacture of cement.

For manufacturing ammonium sulphate fertilizer.


As a drying agent.


Preparation of black board chalks.


Plaster of Paris [CaSO4.1/2H2O]

Calcium sulphate with half a molecule of water per molecule of the salt (hemi-hydrate) is called plaster of paris.

Preparation

It is prepared by heating gypsum (CaSO4.2H2O) at 120°C in rotary kilns, where it gets partially dehydrated.




The temperature should be kept below 140°C otherwise further dehydration will take place and the setting property of the plaster will be partially reduced.

Properties

It is a white powder. When mixed with water (1/3 of its mass), it evolves heat and quickly sets to a hard porous mass within 5 to 15 minutes. During setting, a slight expansion (about 1%) in volume occurs so that it fills the mould completely and takes a sharp impression. The process of setting occurs as follows:




The first step is called the setting stage, and the second, the hardening stage. The setting of Plaster of Paris is catalyzed by sodium chloride, while it is reduced by borax, or alum.

Uses

In surgery for setting broken or fractured bones.


For making casts for statues, in dentistry, for surgical instruments, and toys etc.


In making black board chalks, and statues.


In construction industry.







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Compounds of Aluminium

Alumina

Production
Aluminium oxide, also known as alumina, is the main component of bauxite, the principal ore of aluminium. The largest manufacturers in the world of alumina are Alcoa, Alcan and Rusal.[citation needed] Companies which specialise in the production of speciality aluminium oxides and aluminium hydroxides include Alcan and Almatis. The bauxite ore is made up of impure Al2O3, Fe2O3, and SiO2. Bauxite is purified by the Bayer process:

Al2O3 + 3 H2O + 2 NaOH → 2NaAl(OH)4

The Fe2O3 does not dissolve in the base. The SiO2 dissolves as silicate Si(OH)62-. Upon filtering, Fe2O3 is removed. When the Bayer liquor is cooled, Al(OH)3 precipitates, leaving the silicates in solution. The mixture is then calcined (heated strongly) to give aluminium oxide:[3]

2Al(OH)3 + heat → Al2O3 + 3H2O

The formed Al-2O-3 is alumina.


Uses

Alumina output in 2005Annual world production of alumina is approximately 45 million tonnes, over 90% of which is used in the manufacture of aluminium metal.[3]. The major uses of speciality aluminium oxides are in refractories, ceramics, polishing and abrasive applications. Large tonnages are also used in the manufacture of zeolites, coating titania pigments and as a fire retardant/smoke suppressant.

In lighting and photography, alumina is a medium for chromatography, available in basic (pH 9.5), acidic (pH 4.5 when in water) and neutral formulations. In 1961, GE developed "Lucalox", a transparent alumina used in sodium vapor lamps.[citation needed] Aluminium oxide is also used in preparation of coating suspensions in compact fluorescent lamps.

Health and medical applications include it as a material in hip replacements,[3] in water filters (derived water treatment chemicals such as aluminium sulfate, aluminium chlorohydrate and sodium aluminate, are one of the few methods available to filter water-soluble fluorides out of water), and even in toothpaste formulations.

Aluminium oxide is also used for its strength. Most pre-finished wood flooring now uses aluminium oxide as a hard protective coating. In 2004, 3M developed a technique for making a ceramic composed of aluminium oxide and rare earth elements to produce a strong glass called transparent alumina. Alumina can be grown as a coating on aluminium by anodising or by plasma electrolytic oxidation (see the "Properties" section, above). Both its strength and abrasive characteristics are due to aluminium oxide's great hardness (position 9 on the Mohs scale of mineral hardness).

It is widely used as a coarse or fine abrasive, including as a much less expensive substitute for industrial diamond. Many types of sandpaper use aluminium oxide crystals. In addition, its low heat retention and low specific heat make it widely used in grinding operations, particularly cutoff tools. As the powdery abrasive mineral aloxite, it is a major component, along with silica, of the cue tip "chalk" used in billiards. (See William A. Spinks, cue chalk co-inventor, for more information.) Aluminium oxide powder is used in some CD/DVD polishing and scratch-repair kits. Its polishing qualities are also behind its use in toothpaste.

Aluminium oxide is widely used in the fabrication of superconducting devices, particularly single electron transistors and superconducting quantum interference devices (SQUID), where it is used to form highly resistive quantum tunnelling barriers.

http://en.wikipedia.org/wiki/Aluminium_oxide

Aluminium chloride

Preparation
Aluminium chloride is manufactured on a large scale by the exothermic reaction of aluminium metal with chlorine or hydrogen chloride.[1]

2 Al + 3 Cl2 → 2 AlCl3
2 Al + 6 HCl → 2 AlCl3 + 3 H2

Hydrated forms are prepared by dissolving aluminium oxides with hydrochloric acid.

The solid has a low melting and boiling point, and is covalently bonded. It sublimes at 178 °C. Molten AlCl3 conducts electricity poorly, unlike more ionic halides such as sodium chloride. It exists in the solid state as a six-coordinate layer lattice.



Aluminium chloride is a powerful Lewis acid, capable of forming stable Lewis acid-base adducts with even weak Lewis bases such as benzophenone or mesitylene.[3] Not surprisingly it forms AlCl4− in the presence of chloride ion.

In water, partial hydrolysis forms HCl gas or H3O+, as described in the overview above. Aqueous solutions behave similarly to other aluminium salts containing hydrated Al3+ ions, giving a gelatinous precipitate of aluminium hydroxide upon reaction with the correct quantity of aqueous sodium hydroxide:

AlCl3(aq) + 3 NaOH(aq) → Al(OH)3(s) + 3NaCl(aq)

http://en.wikipedia.org/wiki/Aluminium_chloride

Alum

Alum is a salt that in chemistry is a combination of an alkali metal, such as sodium, potassium, or ammonium and a trivalent metal, such as aluminum, iron, or chromium.

The most common form, potassium aluminum sulfate, or potash alum, is one form that has been used in food processing. Another, sodium aluminum sulfate, is an ingredient in commercially produced baking powder.

The potassium-based alum has been used to produce crisp cucumber and watermelon-rind pickles.

Alum, refers to a specific chemical compound and a class of chemical compounds.
The specific compound is the hydrated aluminum potassium sulfate with the formula KAl(SO4)2.12H2O.

The class of compounds known as alums have the related stoichiometry, AB(SO4)2.12H2O.

Production

Alum from alunite

In order to obtain alum from alunite, it is calcined and then exposed to the action of air for a considerable time. During this exposure it is kept continually moistened with water, so that it ultimately falls to a very fine powder. This powder is then lixiviated with hot water, the liquor decanted, and the alum allowed to crystallize. The alum schists employed in the manufacture of alum are mixtures of iron pyrite, aluminium silicate and various bituminous substances, and are found in upper Bavaria, Bohemia, Belgium, and Scotland. These are either roasted or exposed to the weathering action of the air. In the roasting process, sulfuric acid is formed and acts on the clay to form aluminium sulfate, a similar condition of affairs being produced during weathering. The mass is now systematically extracted with water, and a solution of aluminium sulfate of specific gravity 1.16 is prepared. This solution is allowed to stand for some time (in order that any calcium sulfate and basic ferric sulfate may separate), and is then evaporated until ferrous sulfate crystallizes on cooling; it is then drawn off and evaporated until it attains a specific gravity of 1.40. It is now allowed to stand for some time, decanted from any sediment, and finally mixed with the calculated quantity of potassium sulfa te (or if ammonium alum is required, with ammonium sulfate), well agitated, and the alum is thrown down as a finely-divided precipitate of alum meal. If much iron should be present in the shale then it is preferable to use potassium chloride in place of potassium sulfate.


Alum from clays or bauxite

In the preparation of alum from clays or from bauxite, the material is gently calcined, then mixed with sulfuric acid and heated gradually to boiling; it is allowed to stand for some time, the clear solution drawn off and mixed with acid potassium sulfate and allowed to crystallize. When cryolite is used for the preparation of alum, it is mixed with calcium carbonate and heated. By this means, sodium aluminate is formed; it is then extracted with water and precipitated either by sodium bicarbonate or by passing a current of carbon dioxide through the solution. The precipitate is then dissolved in sulfuric acid, the requisite amount of potassium sulfate added and the solution allowed to crystallize.

Types of alum

Soda alum
Sodium alum, Na2SO4·Al2(SO4)3·24H2O, occurs in nature as the mineral mendozite. It is very soluble in water, and is extremely difficult to purify. In the preparation of this salt, it is preferable to mix the component solutions in the cold, and to evaporate them at a temperature not exceeding 60 °C. 100 parts of water dissolve 110 parts of sodium alum at 0 °C, and 51 parts at 16 °C. Soda alum is used in the acidulent of food as well as in the manufacture of baking powder.


Ammonium alum
Ammonia alum, NH4Al(SO4)2·12H2O, a white crystalline double sulfate of aluminium, is used in water purification, in vegetable glues, in porcelain cements, in natural deodorants (though potassium alum is more commonly used), in tanning, dyeing and in fireproofing textiles.
Uses

Alum in Makeup: Alum was often used as a base in skin whiteners and treatments during the late 16th Century in the Elizabethan fashion.

Shaving alum: is a powdered form of alum used as an astringent to prevent bleeding from small shaving cuts. The styptic pencils sold for this purpose contain aluminium sulfate or potassium aluminium sulfate. Similar products are also used on animals to prevent bleeding after nail-clipping. Alum in block form (usually potassium alum) is used as an aftershave, rubbed over the wet freshly shaved face.

Hair Stiffener: Alum was used in rock form in the 1950's to rub on the front short hair of a "crewcut". When the hair dried, it would stay up all day.
Crystal deodorant: Alum was used in the past as a natural underarm deodorant in Europe, Mexico, Thailand, the Far East and in the Philippines where it is called Tawas. It is now commercially sold for this purpose in many countries, often in a plastic case that protects the crystal and makes it resemble other non-liquid deodorants. Typically potassium alum is used.
Alum powder, found amongst spices at most grocery stores, is used in pickling recipes as a preservative, to maintain crispness, and as an ingredient in some play dough recipes. It is also commonly cited as a home remedy or pain relief for canker sores.
Fire retardant: By soaking and then drying cloth and paper materials they can be made fireproof.
Wax: Alum is used in the Middle East as a component in wax, compounded with other ingredients to create a hair-removal substance.
Foamite: Alum is used to make foamite which is used in many fire extinguishers for chemical and oil fires.
Adjuvant: Alum is used regularly as an adjuvant (enhances immune response to a given immunogen when given with it) in human immunizations.
Antibacterial agent: Alum works as a deodorant because Alum inhibits bacterial growth. This fits the definition of an antibacterial agent. Styptic pencils or Alum powder/crystals can be applied to cuts that have a mild infection.

http://en.wikipedia.org/wiki/Alum
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web sites
thermal decomposition of carbonates, hydroxides and nitrates
http://www.docbrown.info/page01/ExIndChem/ExIndChem.htm-----------------
JEE Question 2007 paper II

Statement - 1

Alkali metals dissolve in liquid ammonia to give blue solutions.

Because

Statement - 2

Alkali metals in liquid ammonia give solvated species of the type [M(NH-3)-n]^+ (M= alkali metals.

(A) Statement – 1 is True, Statement – 2 is True; Statement – 2 is a correct explanation for
statement – 1
(B) Statement – 1 is True, Statement – 2 is True; Statement – 2 is Not a correct explanation for
Statement – 1.
(C) Statement – 1 is True, Statement – 2 is False
(D) Statement – 1 is False, Statement – 2 is True

Solution: B


Blue colour appears due to solvated electrons in liquid ammonia.
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